Resonance booklet of Coordination Compounds

January 9, 2018 | Author: Alok pandey | Category: Coordination Complex, Ligand, Salt (Chemistry), Chelation, Chemical Bond
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LECTURE # 1 1.


Salt : OR OR

The compound which contains acidic(anionic) or basic(cationic) radicals is known as salt. The compound which is formed by the neutralisation reaction of acid and base is known as salt. The compound which on hydrolysis gives acid and base is known as salt.

The most appropriate classification in context of present chapter will be as follows: (1)

Simple salt : The salt which contains one type of simple +ve ion and one type of simple –ve ion. e. g., NaCl, KCl, Na2CO3 etc.


Mixed salt : The salt which contains more than one type of +ve ions or more than one type of -ve ions or both the ions are more than one type. e.g., microcosmic salt NaNH4HPO4.4H2O, bleaching powder Ca(OCl)Cl.


Double salt : When two or more than two simple salts aqueous solutions are mixed in appropriate proportion and then are evaporated, the residue which will be left is known as double salt. when we prepare an aqueous solution of these double salts these will get ionised into all constituent ions. Hence the constituent ions do not loose there identity in double salts.(Generally in double salts one ion is common, either +ve ion or -ve ion in the two or more than two simple salts being mixed). Examples are potash alum - K2SO4 Al2(SO4)3 24H2O, mohr salt - (NH4)2SO4 FeSO4 6H2O and carnallite -KCl MgCl2 6H2O. K2SO4 + Al2(SO4)3 + 24H2O  K2SO4 . Al2(SO4)3 . 24H2O

2K+ (aq.) + 2Al+3 (aq.) + 4SO42– (aq.)


Complex salt/coordination compounds : A complex or a metal complex (sometimes even some electropositve non-metals like Se & Br also form complexes) is defined as a species in which metal atom or ion is bonded by donating some electrons pairs via coordinate bonds and these ions are found to retain their identity in aqueos solutions, hence the original constituent ions loose thier identity and forms a new group which acts as single unit in aqueos solutions (Do remember that this unit/group can be ungrouped or change its identity via chemical reactions). And the complex salts do atleast contain one complex ion. Do not that a complex ion or part is always written in square brackets. Fe(CN)2 + 4KCN  Fe(CN)2 . 4KCN or K4 [Fe(CN)6] (aq.) 4K+ (aq.) + [Fe(CN)6] (aq.) 4:1 comlex Other examples are, [Co(NH3)6]Cl3 (aq) [Co(NH3)6]3+(aq) + 3Cl– (aq) 1: 3 complex [Cu(NH3)4]SO4 (aq.) K2[Zn(CN)4] (aq.)

[Cu(NH3)4]2+ (aq.) + SO42– (aq.) 1 : 1 complex 2K+ (aq.) + [Zn(CN)4]2– (aq.) 2 : 1 complex

Other examples can be, K3[Fe(CN)6], [Cr(H2O)6]Cl3, [Pt(py)4][PtCl4] Coordination compounds are also acid-base adduct and are frequently called complexes or, if charged then complex ion.

(Also note the 1 : 4, 1 : 1, 2 : 1 or 3 : 1 terminology, often used in problems) 2.

DIFFERENT TERMS/DEFINITIONS TO BE USED Some important terms required for the description of coordination compounds are:


COORDINATION ENTITY(COMPLEX) OR COORDINATION SPHERE : A coordination entity constitutes a central atom/ion, usually of a metal, to which are attached a fixed number of other atoms or groups each of which is called a ligand. It may be neutral or charged. In aqueous solutions this complex retain its identity. Examples being : [Co(NH3)6]3+, [PtCl4]2–, [Fe(CN)6]3–, [NiCl2(OH2)4]. The remaining ions apart from complex ions are called Counter Ions or Free Ions or Ionisation Sphere or



Outer Sphere. Ex. Sol.

Designate the coordination entities and counter ions in the coordination compounds : [Cr(NH3)6]Cl3 ; K4[Fe(CN)6], K2[PtCl4] ; [Ni(CO)4]; K2[Ni(CN)4]. The respectively coordination entities are : [Cr(NH3)6]2– ; [Fe(CN)6]4– ; [PtCl4]2– ; [Ni(CO)4]; and the counter ions are Cl– , K+, K+, (no counter ion) and K+, respectively in the given coordination compounds.


CENTRAL ATOM/ION : In a coordination entity – the atom/ion to which are bound a fixed number of ligands in a definite geometrical arrangement around it, is called the central atom or ion. For example, the central atom/ion in the coordination entities : [NiCl2(OH2)4], [CoCl(NH3)5]2+, [Fe(CN)6]3– are Ni2+, Co3+ and Fe3+, respectively.


LIGANDS : The ligands are the ions or molecules bound to the central atom/ion in the coordination entity. This is better visualized as the combination of a Lewis acid (the central atom/ion) with a number of Lewis bases (ligands). The atom of Lewis base that forms the bond to the Lewis acid (central atom/ion) is called donor atom (because it donates the pair of electrons required for bond formation). The central atom/ ion is the acceptor atom/ion (because it receives the electron pairs from the ligands). Some of the common ligands in coordination compounds are : Br–, Cl–, CN–, OH–, O2–, CO32–, NO2–, C2O42–, NH3, CO, H2O, NH2CH2CH2NH2 (1,2-ethanediamine or enthane-1, 2 diamine).

CLASSIFICATION OF LIGANDS (ON BASIS OF DENTICITY) : DENTICITY AND CHELATION : When coordination of more than one sigma-electron pair donor group from the ligand to the same central atom/ion takes place, it is called chelation, and the ligand a chelating ligand. The number of such sigma bonds by such a ligand indicate the denticity of the ligand. For example : unidentate, didentate, terdentate or tridentate, tetradentate, etc. CHELATE LIGAND : Chelate ligand is a di or polydentate ligand which uses its two or more donor atoms to bind a single metal ion producing a ring. The complex formed is referred to as a chelate complex and the process of chelate formation is called chelation. Chelate rings may have any number of atoms; the most common contain five or six atoms, including the metal ion. Smaller rings have angles and distances that lead to strain; larger rings frequently result in crowding both within the ring and between adjoining ligands. Some ligands can form more than one ring; ethylene diaminetetracetate (EDTA) can form five by using the for carboxylate groups and the two amine nitrogens.

The chelate complexes are more stable than similar complexes containing unidentate ligands. The greater stability of the chelate complex in comparison to normal complex is called chelate effect. For example, Ni2+ (aq) + 6 NH3 (aq) [Ni(NH3)6]2+ (aq) Kformation = 108 2+ 2+ Ni (aq) + 3 NH2CH2CH2NH2(aq) [Ni(en)3] (aq) Kformation = 1018 The five and six membered rings are more stable. The following are type and examples if ligands on the basis of this classification. (a) Monodentate (one donor atom) :  : , N  H , CN– , NO – , OH–, CO, NO, C H N (py) etc. F–, Cl–, Bl–, H O 2





(b) Didentate (two donor atoms : Some most common bidentate ligands are


NH2 – CH2 – CH2 – NH2 Ethylenediamine(en)



NH2  CH  CH2  NH2 (1, 2 propanediamine) (pn) | CH3



C C2O42–


O O Oxalate ion (ox2-) O CO32–



C || O Carbonate ion

2, 2 bipyridyl (bipy)



glycinate ion (gly)

OR acetylacetonate ion (acac) ion

CH3 – C  N – O – | CH3 – C  N – OH


two nitrogen atoms as donor atoms

dimethyl glyoximate ion (dmg)

DIDENTATE CHELATION : In [PtCl2(en)], en represents the didentate ligand, NH2CH2CH2NH2 (1,2-ethanediamine or ethylenediamine) [Figure] H2 N

Cl Pt Cl



CH2 N H2 [PtCl2(en)] (a)


(c) TRIDENTATE/TERDENTATE LIGAND/CHELATION :  H – CH – CH – N  – CH – CH – N  H H2N 2 2 2 2 2 Diethylenetriamine (dien)


[N-(2-aminoethyl)-1, 2-ethanediamine]

TRIDENTATE/TERDENTATE : In the coordination entity (Pt Cl(dien)]+, dien, [N-(2-aminoethyl)-1, 2-ethanediamine] is a terdentate ligand [Figure] CH2


H2N Pt



N H2




[PtCl(dien)] (b)

(d) TETRADENTATE LIGAND/CHELATION :  H – CH – CH –  – CH – CH – N  – CH – CH – N  H H2N NH 2 2 2 2 2 2 2 Triethylenetetraamine (trien) OR [N, N+-bis-(2 aminoethyl)-1, 2-ethanediamine] In [Pt(trien)2+], trien, [N, N+-bis-(2 aminoethyl)-1, 2-ethanediamine represents a tetradentate ligand. H2C – CH2 H N






N H2


N H2 [Pt (trien)] (c)



 H – CH – CH – N  H  H – CH – CH – N  H – CH – CH – N  – CH – CH – N H2N 2 2 2 2 2 2 2 2 2 tetraethylenepentaamine Ethelene diamine triacetate ion. O



– O–C–CH – O–C–CH


N–(CH2 )2–N

CH2–C–O – H



O (Ethelenediaminetriacetato ion)

(f) HEXADENTATE LIGAND/CHELATION : EDTA is the most common hexadentate ligand and it is derived from ethylene diamine tetraacetic acid

This acid is directly not soluble in water but its sodium salt is soluble. O C


CH2 C :N







: :

: :



: :

: :







O EDTA4–, Ethylenediaminetetraacetate ion

Therefore, acid is not the ligand but its tetra anion is a ligand and its usual denticity is 6 but it is sometimes also found to behave as tetradentate ligand and hence is a flexidentate ligand.



(g) FLEXIDENTATE LIGAND : A polydentate ligand which is found to have different denticity in different coordination compounds is called a flexidentate ligand. Note that in a particular complex denticity of a particular ligand is fixed, it can not be flexible in the same compound. EDTA can act as hexa, penta as well as tetra dentate ligand. For example ; EDTA usually acts as hexadentate ligand but in [Cr(III)(OH)(EDTA)]2– and [Co(III)Br(EDTA)]2– as pentadentate and in [Pd(II)H2(EDTA)]0 as a tetradentate ligand. Sulphate ion, SO42– can also be mono or bi dentate ligand. For example ;

(h) AMBIDENTATE LIGAND : Ligands which can ligate through two different atoms present in it are called ambidentate ligands. Examples of such ligands are the NO2– and SCN¯ ions. NO2– ion can coordinate through either the nitrogen or the oxygen atoms to a central metal atom/ion. Similarly, SCN¯ ion can coordinate through the sulphur or nitrogen atom. Such possibilities give rise to linkage isomerism in coordination compounds. For examples : CN–




– bonding may be through through C or N – CN , – NC – bonding may be through N or O.


M  O—N=O M  SCN M  NCS


nitrito-O thiocyanato or thiocyanato-S isothiocyanato or thiocyanato-N

OXIDATION NUMBER OF CENTRAL ATOM : The oxidation number of the central atom is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom. Oxidation number is represented by a Roman numeral in parenthesis following the name of the coordination entity. Some examplex are listed Table below.

Examples of important terms used in describing coordination entities : –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– Coordination entity (complex) Ligand list Central atom/oxidation number –––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––– [Co(NH3)6]3+ 6NH3 Co/ (III) [NiCl4]2– 4Cl– Ni / (II) [Co(CN)5F–]3– 5CN– + 1F– Co / (III) [Ni(CN)4]2– 4CN¯ Ni / (II) [Ni(CO)4 4CO Ni / (O) [Ni(H2O)6]2+ 6H2O Ni / (II) ––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––––



LECTURE # 2 3.

NOMENCLATURE OF COORDINATION COMPOUNDS : We first are writing down the rules to write down the names of two impotant parts of a complex, first of ligands and then that of central atom. Naming of Ligands: The names of ligands are written according to the charge (+ , – or nuetral) on them. Naming of Anionic ligands: the name of these ligands ends with –o. The name of anions can be of following three types 1. If name of the anion ends with –ide then generally –e is replaced with –o 2. If the name of anion ends with –ate then –e is replaced with –o 3. If the name of anion ends with –ite then –e is replaced with –o H– hydrido O2– oxido OH– hydroxido – X halido(chlorido, flourido, bromido, iodido) O22– peroxido O2– superoxido O3– ozonido N3– nitrido N3– azido NH2– amido NH2– imido CN– NC– [NCO]– [SCN]– NO2–

CH3COO– NO3– S2O32– SO4– SO3– HSO3– HO2– (NOS)– S2– HS– H2N – CH2COO–

cyanido or cyanido–C (second is preferred) isocyanido or cyanido–N cyanato –O or –N thiocyanato–N or –S nitro(is common when donation is through nitrogen) –NO2 nitrito–N –ONO nitrito–O acetato nitrato thiosulphato sulphato sulphito hydrogensulphito hydrogenperoxido thionitrito thio or sulphido hydrogensulphido(mercapto) RSH  mercatanes (greater affinity to mercury) glycinato (gly) acetylacetonato (acac)

CH3 – C = N – O– | CH3 – C = N – OH

dimethylglyoximato (dmg)

Naming of Neutral ligands : Generally names of neutral ligands are not modified except in some cases some special names are used for some common neutral ligands. NH3 ammine H2O aqua or aquo S8 octasulphur P4 tetraphosphorus CH3 NH2 methylamine



C5 H5 N NH2 – CH2 – CH2 – NH2 O2 CO NO

pyridine (py) ethylenediamine (en) or ethane -1, 2-diamine dioxygen carbonyl nitrosyl

o - phenanthroline (o-phen) N..



2, 2 bipyridyl (bipy) C2H4 C6H6

ethylene or ethene benzene

NOTE : IN LAST TWO CASES THE PI ELECTRIN DENSITY IS GENERALLY DONATED Positive/Cationic ligands : The name of these ligands ends with –ium NO+ N2H5+ NO2+

: : :

nitrosonium or nitrosylium (old) hydrazinium nitronium

Naming of central atom/ion The name of central atom is not modified if the central atom is in the cationic part of the complex while its name is modified and it ends with –ate if it is present in the anionic part of the complex. Metal Name in Cationic part Name in Anionic part Cr chromium chromate Pt platinium platinate Co cobalt cobaltate Ni nickel nickelate Zn zinc zincate Pd palladium palladate Ti titanium titanate Va vanadium vanadate Mo molybdenum molybdate Hg mercury mercurate W tungstun tungstate Si silicon silicate Br bromine bromate B boron borate Os osmium osmate Cu copper cuprate Note : For some metals, the Latin names are used in the complex anions. Pb lead plumbate Ag silver argentate Au gold aurate Sn tin stannate Fe iron ferrate Note : The name of central atom is not modified if the central atom is in the neutral part of the complex.



Steps for Writing the Name of a Complex :

(i) Calculate oxidation number of the central atom involved. (ii) The cation is named first as is the case with other ionic compounds. This rule applies to both positively and negatively charged coordination entities. For example in K4[Fe(CN)6] and [Co(NH3)6]Cl3, K+ and [Co(NH3)6]3+ cations are named first. Potassium hexacynidoferrate(II) Hexaaminecobalt(III) chloride (iii) Usually the number of free cation or anions is not given as can easily be calculated by charge balancing. K4 [Fe(CN)6] [Co(NH3)6] Cl3 Potassium but not tetrapotassium Chloride but not trichloride (iv) For writing the names of complex cation use the following rule. Names of ligand in english alphabetical order with their qunatity as prefix (but do not use mono for one and do not consider first letters of these number denoting prefixes in alphabetisation). This is followed by the name of metal followed by its oxidation number (in Roman numeral) in a brackette like (I), (II), (III), (O), (–I), (–II), etc. [Co(NH3)4Cl2]+, pentaamminechloridocobalt(III). (v) For writing the names of complex anion use the following rule. Names of ligand in english alphabetical order with their qunatity as prefix (but do not use mono for one and do not consider first letters of these number denoting prefixes in alphabetisation). This is followed by the name of metal(modified as per rule already specified) followed by its oxidation number (in Roman numeral) in a brackette. (I), (II), (III), (O), (–I), (–II), etc. (NH4)2 [Co(SCN)4], ammonium tetrathiocyanato-S-cobaltate(II). Note : When the names of the ligands include a numerical prefix or are complicated or whenever the use of normal prefixes creates some confusion, it is set off in parentheses and the second set of prefixes is used. 2 di bis 3 tri tris 4 tetra tetrakis 5 penta pentakis 6 hexa hexakis 7 hepta heptakis [CoCl2(NH2CH2CH2NH2)2]+, dichloridobis(ethane-1,2-diamine)cobalt(III). [NiCl2(PPh3)2], dichloridobis(triphenylphosphine)nickel(II). (vi) The neutral complex molecule is named similar to that of the complex cation. Examples ;

[Pt(NH3)BrCl(CH3NH2)], amminebromidochloridomethylamineplatinum(II). EXAMPLES ILLUSTRATING THE USE OF RULES OF NOMENCLATURE FOR COORDINATION COMPOUNDS: Formula of Coordination Compound K3[Fe(CN)6]

Name of Coordination Compound Potassium hexacyanidoferrate(III)


Pentaamminechloridocobalt(III) chloride


Diamminechlorido(methylamine)platinum(II) chloride


Potassium tetrachloridopalladate(II)


Tetraamminediaquacobalt(III) chloride




Tetraamminediaquanickel(II) sulphate


Potassium tetracyanidonickelate(II)


Hexaamminecobalt(III) chloride sulpahte


Iron(III) hexacyanidoferrate(II)




Sodium amminebromidochloridonitrito–N–platinate(II)


Potassium trioxalatoaluminate(III) or Potassium tris(oxalato)aluminate(III)


Hexamminecobalt(III) diamminetetranitrito-N-cobaltate(III)


Hexakis(methylcyanide)iron(II) Bromide

[CuCl2 {O = C (NH2)2}2]


RULES FOR WRITING THE FORMULAE OF MONONUCLEAR COORDINATION COMPOUNDS : Mononuclear coordination entities contain a single central metal atom. The sequence of symbols within the formula of coordination entity is governed by the following rules : (i) The central atom is listed first ; (ii) The ligands are then placed in alphabetical order. The placement of a ligand in the list does not depend on its charge. (iii) Polydentate ligands are also placed alphabetically. In case of abbreviated ligand, the first letter of the abbreviation is used to determine the position of the ligand in the alphabetical order. When ligands are polyatomic, their formulas are enclosed in parentheses. Ligands abbreviations are also enclosed in parentheses. (iv) The formula of the coordination entity is enclosed in square brackets. (v) No space is kept between representations of ionic species within the formula ; (vi) When the formula of a charged coordination entity is written without the formula of the counter ion, the charge is indicated outside the square brackets as a right superscript with the number before the sign. (vii) The charge of the cation(s) is balanced by the charge of the anion(s). For example : [Co(CN)6]3–, [Cr(H2O)6]3+ etc. The following examples illustrate the above rules : [Co(NH3)6]Cl3 ; [Co(NH3)5Cl]Cl2 ; K2[PtCl4] [Co(NH3)4Cl(NO2)]Cl ; K3[Fe(CN)6] ; Na[AllII(H2O)2(OH)4] Ex.


Write the formulae for the following coordination compounds : (i) tetrahydroxidozincate (II) (ii) pentaaquachloridochromium (III) chloride (iii) tetrabromidocuparate (II) (iv) pentacarbonyliron (O) (v) potassium tetracyanidocuprate (II) (i) Zn(OH)4]2– (ii) [Cr(OH2)5Cl]Cl2 (iii) [CuBr4]2– (iv) [Fe(CO)5]

(v) K2[Cu(CN)4]

Naming of Some Special Complexes and of Complexes with Bridging ligands : The name of a bridge complex is prefixed by – If the situation on both sides of the bridge is symmetrical then we can write the name of remainig complex at one place like (i)

 3 1  3   (NH3 )5 Cr  OH – Cr (NH3 )5   



–hydroxidobis(pentaamminechromium(III)) chloride OR we could also have written the name of one side of the bridge then the name of bridging complex and then the name of the other side of the bridge, like  3 1  3   (NH3 )5 Cr  OH – Cr (NH3 )5   



pentaamminechromium(III)––hydroxidopentaamminechromium(III) chloride But if the compound is unsymmetrical on both sides of the bridge then we have to follow the second rule , i.e. write the name of one side then that of the bridge and then that of the second side, like

Tetraaquacobalt(III)––amido––hydroxidotetramminechromium(III) sulphate Some other examples are



(ii) –amido––hydroxidobis(tetramminecobalt (III)) ion


(SO4)2 di––hydroxidobis(tetraaquairon(III)) sulphate


(SO4)2 –amido––hydroxidobis(bis(ethylenediammine)cobalt(III)) sulphate


bis(ethylenediammine)cobalt(III)––amido––hydroxidobis(ethylenediammine)cobalt(III) sulphate.

Naming of special compounds :


The name of the above compund is bis(cyclopentadienyl)iron(II). This is commonly known as Ferrocene. In the above compound the aromatic pi electron density of the ring is being donated to the Fe2+ ion, hence in this way iron is bonded to all the five carbons of the upper flat ring and also to all five carbons of below flat ring. This is also called a sandwich compound because of its shape. Its not a true complex this is an organometallic compound which we study in detail in the last part of the chapter. To write down the formula of such compound a special symbol  is written before the name of such ligand and on superscript to this  the number of atoms is written between which the electron density is being shared. So the formula of this compound will be Fe[  5–C5H5]2 . The other such example is bis(benzene)chromium which will be represented by the formula Cr[  6–C6H6]2 – H





Pt H Cl

H Cl

The last compound is commonly known as Zeisse’s salt. The name of the compound will be trichloridoethyleneplatinate(II) ion and the formula will be written as [PtCl3(  2–C2H4)]



LECTURE # 3 (i) (a)

BONDING IN COORDINATION COMPOUNDS : Initial bonding theories and EAN rule WERNER’S COORDINATION THEORY The systematic study of coordination compounds was started by Alfred Werner whose pioneering work opened an entirely new field of investigation in inorganic chemistry. He prepared and characterized a large number of coordination compounds and studied their physical, chemical and isomeric behaviour by simple experimental techniques. On the basis of these studies. Werner, in 1898, propounded his theory of coordination compounds.

The main postulates of Werner’s theory are : 

Metals exert two types of linkages/valencies ; (i) the primary or ionizable links/valencies which are satisfied by negative ions and equal the oxidation state of the metal and (ii) the Secondary or nonionizable links/ valencies which can be satisfied by neutral or negative ions/groups(or sometimes by cationic species discovered later). The secondary linkages equal the coordination number of central metal atom/ion. This number is fixed for a metal.

The ions/groups bound by the secondary linkages have characteristic spatial arrangements corresponding to different co-ordination numbers. In the modern terminology, such spatial arrangements are called coordination polyhedra. The various possibilities are CN = 2 Linear CN = 3 Triangular CN = 4 Tetrahedral or sq. planar

CN = 5

square pyramidal or TBP

CN = 6


The secondary valencies are generally represented by solid lines while the primary valencies are represented by dashed lines and the ions which satisfy both primary and secondary valencies will be drawn with both solid and dashed lines. For example the complex [CoCl (H2O)5] Cl2 is represented as

On the basis of the above postulates, Werner formulated the coordination compounds, CoCl3 . 6NH3, CoCl3 . 5NH3 and CoCl3 . 5NH3 and CoCl3 . 4NH3 as : [Co(NH3)6]Cl3, [CoCl(NH3)5]Cl2 and [CoCl2(NH3)4]Cl. respectively - the species within the square brackets being the coordination entitles (complexes) and the ions outside the square brackets the counter ions. He further postulated that octahedral, square, planar and tetrahedral geometrical shapes are more common in coordination compounds of transition metals. Thus, [Co(NH3)6]3+, [CoCl(NH3)5]2+, [CoCl2(NH3)4]+ are octahedral entities, while [Ni(CO)4] and [PtCl4]2– are tetrahedral and square-planar, respectively. Series of coloured compounds obtained by the interaction of aqueous CoCl3 and NH3 Compound Colour Name according to colour CoCl3 . 6NH3 Yellow Luteo Complex CoCl3 . 5NH3 Purple Purpureo Complex CoCl3 . 4NH3 Green Praseo Complex CoCl3 . 4NH3 Violet Violeo Complex



Werner used some experimental methods to find out the number of free ions and coordination number in many different complexes. Such methods involved freezing point dpression measurement or boiling point elevation measurement of aqueous solution of these complexes( we will study later about these methods). Other methods included conductivity measurement of the aqueous solutions of these complexes. Assume here that the conductivity of a solution only depends on the number of ions present per unit volume or on concentration of the ions( which actually is not the case, we will study later that it also depends on the nature of the on), try and arrange the solution of the following complexes in increasing order of conductivity. S.No.

Werner complex

Modern notation





[Co(NH3)6] +3Cl




[Co(NH3)5Cl] +2Cl








Secondary valency satisfied by

Primary valency satisfied by

six (NH3)

three (Cl )


five (NH3) and one (Cl )

[Co(NH3)4Cl2] + Cl


four (NH3) and two (Cl )


three (NH3) and three (Cl )

three (Cl ) including one (Cl ) with dual nature – three (Cl ) including – two (Cl ) with dual nature –

three (Cl ) all with dual nature

From the above table it is clear that (i) the solution conductivity of complexes 1,2 and 3 corresponds to 1:3, 1:2 and 1:1 electrolyte respectively and thus the increasing order of the conductivity can be represented as CoCl33NH3 < CoCl3. 4NH3 < CoCl3.5NH3 < CoCl3.6NH3 and (ii) the complexes 1,2 and 3 will react with silver nitrate and give 3, 2 and 1 mole of the white precipitate of silver chloride. Ex. Sol.


Ans. Sol.

For the compound represented by the formula CoCl3.5H2O write down the different complexes possible and repesent them as per Werners notation and also write down their names. CN = 6 octahedral [CoCl(H2O)5]Cl2; [CoCl2(H2O)4]Cl.H2O; [CoCl3(H2O)3].2H2O; generally water of crystallization is found with complexes which have bogth cationic and anionic parts, it is less common with species which are neutral.

Match the pairs of complexes listed in column-I with the method(s) used for their differentiation listed in column-II. Column – I Column–II (A) [Cr(H2O)6]Cl3 and Cr(H2O)5Cl] Cl2.H2O

(p) Can be differentiated by amount, nature or colour of precipitate formed.

(B) [Co(NH3)5Br]SO4 and [Co(NH3)5SO4]Br

(q) Can be differentiated by electrical conduction measurement method (appreciable difference)

(C) [Co(NH3)5Cl]Cl2 and [Co(NH3)6]Cl3

(r) Can be differentiated using cryoscopic measurement method.

(D) [Cu(H2O)4]SO4.H2O and [Cu(H2O)6](NO3)2

(s) Can be differentiated by heating with concentrated H2SO4

(A) – p,q,r,s ; (B) – p,q ; (C) – p,q,r ; (D) – p,r,s (A) [Cr(H2O)6]Cl3 (aq) [Cr(H2O)6]3+ (aq) + 3Cl– (aq) [Cr(H2O)5Cl] Cl2.H2O (aq) (B)


[Co(NH3)5Br]SO4 (aq)

[Co(NH3)5Br]2+ (aq) + SO42– (aq)

[Co(NH3)5SO4]Br (aq)

[Co(NH3)5SO4]+ (aq) + Br– (aq)

[Co(NH3)5Cl]Cl2 (aq) [Co(NH3)6]Cl3 (aq)


[Cr(H2O)5Cl]2+ (aq) + 2Cl– (aq)

[Cu(H2O)4]SO4.H2O (aq) [Cu(H2O)6](NO3)2 (aq)


[Co(NH3)5Cl]2+ (aq) + 2Cl– (aq) [Co(NH3)6]3+ (aq) + 3Cl– (aq) [Cu(H2O)4]2+ (aq) + SO42– (aq) [Cu(H2O)6]2+ (aq) + 2NO3– (aq) COORDINATION COMPOUND - 13

Note : (1) (2) (3) (4) (b)

Ag+ + Cl–  AgCl  (white)

; Ag+ + Br–  AgBr  (pale yellow)

Ba2+ + SO42–  BaSO4  (white) Conductivity depends on ; (a) the charge on cation and anion and (b) the number of ions in aqueous solution. Cryoscopic measurement method depends on the number of ions in aqueous solution. H2SO4 acts as dehydrating agent and thus removes the water molecules present in the ionisation sphere.

SIDGWICK’S EFFECTIVE ATOMIC NUMBER RULE (EAN) : On the basis of some limited number of complexes available at that time, Sidgwick observed some stable complexes and formulated the following rule, which is found to be incorrect when many more omplexes have been prepared.( So the rule is not scientific, but sometimes questions are directly asked on calculation of EAN, so we should be knowing it.) EAN Rule: The complexes in which the EAN of the central atom equals the atomic number of the next noble gas, are found to be extra stable. EAN(Effective Atomic Number) = No. of e– present on the metal atom/ion + no. of e–s donated by ligands to it. OR

EAN = (Atomic no. of C. M. – O. S. of CM) + [No. of electron donated by ligands] For fast calcultions remember the atomic number of the noble gases which are He 2 Ne 10 Ar 18 Kr 36 for 3d series metals Xe 54 for 4d series metals Rn 86 for 5d series metals Ex.(a) [Co3+(NH3)6]Cl3 Number of electrons in Co3+ = 27 – 3 = 24 6NH3 = 6 × 2 = 12 So, EAN = 24 + 12 = 36 (krypton) (b)

[Fe2+(CN)6]4– Number of electrons in Fe2+ = 26 – 2 = 24 6CN– = 6 × 2 = 12 So, EAN = 24 + 12 = 36 (krypton)


[Fe3+(CN)6]3– Number of electrons in Fe3+ = 26 – 3 = 23 6CN– = 6 × 2 = 12 So, EAN = 23 + 12 = 35

NOTE : (a) The EAN rule is generally found to be not valid in case of most of the complexes but in case of metal carbonyls( which is an important class of comlexes, we will be studying later) this rule is found to be valid in all cases except one or two exceptions, so do remember it for metal carbonyls and do know how to calculate the EAN for any metal. (b)

(i) (ii) (iii)

NO ligand is found to act as three electron donar, as indicated by the following reactions in which when the carbonyl compound is heated in atmosphere of excess of NO. [Fe(CO)5] + 2NO  [Fe(CO)2(NO)2] + 3CO [Cr(CO)6] + 4 NO  [Cr(NO)4] + 6 CO But also remember that NO+ is only two electron donar The Werner theory and some other initial bonding theories cannot answer basic questions like. Why only certain elements possess the remarkable property of forming coordination compounds ? Why the bonds in coordination compounds have directional properties ? Why coordination compounds have characteristic magnetic and optical properties ?



LECTURE # 4 VALENCE BOND THEORY : The valence bond theory, VBT, was extended to coordination compounds by Linus Pauling in 1931. The basic principles involved in this treatment are : (a)

Metal will provide vacant orbitals depending on its CN. These vacant orbitals will be hybridised. Thus hybrid vacant orbital of metal will overlap with the filled orbital of the ligand to form a coordinate bond.


There is also found some kind of back bonding from metals to the ligands to maintain the charge neutrality. The extent of this back bonding also depends on the nature of the ligand. We will study about this a bit in metal carbonyls and all other complexes generally neglect this.

 bond  It is a partial bond between filled orbitals of metal & vacant of ligand. The ligands which allow back bonding to sufficient extent are also called  acid acceptors. For example CO, CN–, NO etc. (c)

The type of hybridisation of metal & shape of complex involved can be decided conveniently if some chracteristics of the complex like magnetic nature, geometry or whether exhibits isomerism or not etc. be known. Hence along with the coordination number some other property must also be given to analyse the complex using VBT. sp3, dsp2, dsp3 and d2sp3 yield tetrahedral, square planar, trigonal-bipyramidal or square-pyramidal and octahedral spatial arrangements commonly encountered in coordination compounds. It is found that [Ni(CN)4]2–, is diamagnetic in nature. But we know that Ni2+ is paramagnetic with two unpaired electrons, hence to be diamagnetic the electrons must have got paired up and one d orbital will be available for the dsp2 hybridisation. For one d-orbital lobe available for hybridisation, pairing of electrons takes place in the remaining d-orbitals, as shown below in the case of [Ni(CN)4]2–, 3d 4s 4p Ni       Ni2+      [Ni(CN)4]2–         CN¯ CN¯ CN¯ CN¯ dsp2 hybrids used [Ni(CN)4]2– is square planar and diamagnetic. [NiCl4]2–, on the other hand, is paramagnetic and has tetrahedral geometry. In this case, the VB treatment assumes that (i) the d-orbital occupancy remains the same as in the free Ni2+ ion and (ii) the metal uses sp3 hybrids (involving 4s and 4p orbitals) for bonding with the ligands as shown below : [NiCl4]2– is paramagnetic as there are two unpaired electrons. 3d 4s 4p Ni       Ni2+      [Ni(CN)4]2–          Cl¯ Cl¯ Cl¯ Cl¯ 3 sp hybrids used The VB theory explains the formation of 6– coordinate, octahdral coordination entities by invoking the use of (n-1)d2nsnp3 or nsnp3nd2 hybrid orbitals by the central metal ion in forming bonds with the ligands. Whether the hybridisation is dsp3 or sp3d, it can be decided if number of unpaired e– or magnetic moment is known to us else not. An example involving use of (n-1)d2nsnp3 hybrids is the coordination entity, [Fe(CN)6]4–, which is founf to be diamagnetic The double occupancy of (n-1)d orbitals orbitals confers extra stability and the absence of unpaired electrons renders this entity diamagnetic.



ns np            L L L L L L (d2sp3 hybrid orbitals filled by electron pairs donated by the respectively ligands. L = CN¯).

Fe Fe2+ [Fe(CN)6]4–

(n-1)d   

  

  

Ex. Sol.

Predict the number of unpaired electrons in the square planar [Pt(CN)4]2– ion. Pt2+ is a 5d8 ion. For square-planar geometry, dsp2 hybrids are required. For the availability of one d orbital, pairing of electrons takes place in the remaining d-orbitals. Hence there are no unpaired electrons in [Pt(CN)4]2– ion.

Ex. Sol.

For the comlex [Ma2 b2], It is given that it does not show geometrical isomerism. Predict the hybridisation. CN = 4 Hence either square planar( dsp2 ) or tetrahedral( sp3 ) but since no geometrical isomerism, hence must be tetrahedral( sp3 ).

Some More Examples With Different Coordination number (CN) : CN = 2 The coordination two is relatively rare, occurring mainly with the +1 cations of Cu, Ag, Au and with Hg2+. The coordination geometry is linear. For example Ex.


Ag()  4d10 5s0

Ag – C bond in [Ag(CN)2]– is obtained by overlapping of sp hybridised orbitals of Ag() and sp hybridised orbitals of C of CN– Such complexes are typically unstable towards the addition of further ligands as [Cu(CN)2]– + 2CN–  [Cu(CN)4]3– CN = 3 The most important geometries for the complexes with CN = 3 are trigonal planar and the trigonal pyramidal. Examples are the planar HgI3– and [Cu(CN)3]2– and the pyramidal SnCl3–. Ex.


Hg()  5d10 6s0 p0

triiodomercurrate (II) ion

C.N. = 4 This is one of the most important coordination number and results in tetrahedral and square planar geometry. Ex.

Ni (CO)4 (liquid), is found to be diamagnetic, explain the shape and hybridisation.



[NiCl4]2–, is found to be paramagnetic with two unpaired electrons, explain the shape and the hybridisation.





Cl– are weak ligand( will study later in CFT) So, they cannot force 3d8 electrons to pair up in the 4 orbitals. So, it is paramagentic complex. Ex. Sol.

[Ni2+(CN)4]2– , the complex is found to be diamagnetic, explain the shape and hybridisation of the complex. 3d8 4s0 p0 d0

square planar


net = 0

CN– are strong field ligands (will study in CFT) and therefore they will force the 3d8 electrons to pair up in 4 orbitals. CN = 5 Coordination number five is less common than coordination number four or six, the possible geometries are trigonal bipyramidal and the square pyramidal. In most of the complexes these geometries are also found to interchange into each other. For both these geometries sp3d and dsp3 are two possible hybridisations. Ex.

[Fe(CO)5], this complex is founf to have its total dipole moment to be equal to zero. Explain the hybridisation.


Fe  3d6 4s2 p0 d0

Trigonal bipyramidal and diamagnetic

net = 0


[Ni(CN)5]3– is found to be dimagnetic with 2 types of Ni – C bond lengths out of which four bond lengths are found to be equal and the fifth one is different. Explain the hybridisation, shape and thus the observed properties.


CN = 5



CN = 6 This is the most common coordination number and the geometry associated is octahedron. For hybridisation there are two possibilities sp3d2 or d2sp3

Outer d complexes Inner d complexes (Octahedral) (Octahedral) High spin = no. of unpaired e–  generally Outer d complexes Low spin = no. of unpaired e–  generally Inner d complexes Ex. Sol.

[Co (NH3)6]3+, this complex is found to be diamagnetic, explain shape and hybridisation. Co3+  3d6 4s0 p0 d0 Therefore, pairing of electrons in 3d sub shell of Co(III) is forced for the complex to be diamagnetic.

Spin = 0 Ex. Sol.

Low spin  NA

[Fe(CN)6]4– complex is found to be diamagnetic while the [Fe(CN)6]3– complex is found to be paramagnetic with one unpaired electron. Explain the shape and hybridisations of both. [Fe(CN)6]4– Fe2+  3d6

diamagentic and low spin complex

[Fe3+(CN)6]3– Fe3+  3d5 Ex. Sol.

paramagentic (I) low spin complex

[FeF6]3– complex is found to be highly paramagnetic with five unpaired electrons, explain the shape and hybridisation. Fe3+  3d5

paramagentic (5) high spin complex/outer -d- complex While the VB theory, to a large extent, explains the formation, structures and magnetic behaviour of coordination compounds, it suffers from the following shortcomings : 

A number of assumptions are involved.

There is no quantitative interpretation of magnetic data.

It has nothing to say about the spectral(color) properties of coordination compounds.

It does not give a quantitative interpretation of the thermodynamic or kinetic stabilities of coordination compounds.

It does not make exact predictions regarding the tetrahedral and square-planar structures of 4-coordinate complexes.

It does not distinguish between strong and weak ligands. The drawbacks of VBT of coordination compounds are, to a considerable extent, removed by the Crystal Field Theory.




CRYSTAL FIELD THEORY : The Crystal Field Theory (CFT) was originally proposed for explaining the optical properties of crystalline solids.It was applied to the study of coordination compounds in the 1950s. CFT assumes the ligands to be point charges and the interaction between them and the electrons of the central metal to be electrostatic in nature. The five d-orbitals in an isolated gaseous metal atom/ion have same energy, i.e., they are degenerate. This degeneracy is maintained if a spherically symmetrical field of negative charges surrounds the metal atom/ion. However, when this negative field is due to ligands (either anions or the negative ends of dipolar molecules like NH3 and H2O) in a complex, it becomes asymmetrical and the degeneracy of the d-orbitals is lifted. It results in splitting of the d-orbital energies. The pattern of splitting depends upon the nature of the crystal field. We will first consider :



dxy dyz dzx planar d orbitals



2 y 2 z2 x  

axial d orbitals

or t2g



For convenience, let us assume that the six ligands are positioned symmetrically along the cartesian axes, with the metal atom at the origin. As the ligands approach, first there is an increase in the energy of d orbitals to that of the free ion just as would be the case in a spherical filed. Next, the orbitals lying along the axes (dz2 and dx 2  y 2 ) get repelled more strongly than dxy, dyz and dxz orbitals, which have lobes directed between the axes. The d xy , d yz , dxz orbitals are lowered in enrgy relative to the average energy in the spherical crystal filed. Thus, the degenerate set of d orbitals get split into two sets : the lower energy orbitals set, t2g and the higher enrgy, eg set. The energy seperation is denoted by 0 (the subscript o is for octahedral (Fig.)



Let us now learn about the significance od 0 by considering first the d1 coordination entity, [Ti(H2O)6]3+, formed in aqueous solutions of Ti3+ (d1) ion. Obviously, the single d electron occupies one of the lower energy t2g orbitals. In d2 and d3 coordiantion entities, the d electrons occupy the t2g orbitals singly in accordance with the Hund’s rule. For d4 ions, two possible patterns of electron distribution arise : i) the fourth electron may enter an eg orbital (of higher energy) or ii) it may pair an electron in the t2g level. The actual configuraion adopted is decided by the relative values of 0 and P. P represents the enrgy required for electron pairing in a single orbital. If is less than P (0 < P), we have the so called weak field, high spin situaiton, and the fourth electron 3 1 entres one of the eg orbitals giving the configuraiton t 2g eg . If now a fifth electron is added to a weak field 3


coordination entity, the configuraiton becoms t 2g eg . When 0 > P, we have the strong field, low spin situation, and pairing will occur in the t2g level with the eg level remaining unoccuped in tentities of d1 to d6 ions. Calculations show that coordination entities with four to seven d electrons are more stable for strong field as compared to weak field cases. This extra stabilization due to ligand field in comparison to normal situation( if there would have been no spilitting or if there would have been spherically symmetrical ligand field) is called CFSE(Crystal field stabilization energy). This CFSE can be calculated as follows

General formula : CFSE = [– 0.4 (n) t2g + 0.6 (n) eg] o + nP. where n & n are number of electron(s) in t2g & eg orbitals respectively and o crystal field splitting energy for octahedral complex. n represents the number of extra electron pairs formed because of the ligands in comparison to normal degenerate configuration. Metal ion Configuration


Configuration in Ligand field



Sc 3+

t2g0,0,0, eg0




t2g1,0,0, eg0

–0.4 0










Cr , V 2+




1,1,1 2,2,2

, eg


, eg



–0.8 0 –1.2 0


–2.4 0 + 1.2 0 = –1.2 0






t2g2,2,2 eg2,1

–2.4 0 + 1.8 0 = –0.6 0



t2g2,2,2 eg2,2

–2.4 0 + 2.4 0 = 0

Therefore, for the above configurations, there is no effect of the nature of ligand. They may be strong or weak, the formula for CFSE will remain the same. For configuration from d4 to d7 : Metal ion Configuration 4


Example 2+

Cr S.L. W.L.

` 5



Mn S.L.

Conf. in Ligand field








, eg


, eg


, eg


–1.6 0 + P –0.6 0 –2.0 0 + 2P


d6 7


Fe W.L.

t2g1,1,1, eg1,1

Co3+,Fe2+ S.L.

t2g2,2,2, eg0,0

–2.4 0 + 2P


t2g2,1,1, eg1,1

– 0.4 0


Co S.L. W.L.






, eg


, eg


–1.8 0 + P – 0.8 0

Note : S.L. = Strong field ligands W.L. = Weak field ligands Ex.

For the complex[Cr(H2O)6]2+ , P = 23500 cm–1 and its hybridisation and net CFSE in kJ/mol.


Cr2+  3d4  t2g1, 1, 1 eg1


P > 0 weak ligand


0 = 12600 cm–1, then for this complex calculate

outerd complex.

CFSE = (– 3 × 0.4 + 1 × 0.6) 0 = – 0.6 0 net stabilisation energy = CFSE + P (wherever there is pairing)


= CFSE + 0

= – 0.6 0


SPECTROCHEMICAL SERIES On the basis of the behaviour (the splitting produced with different metal ions) of ligands these are arranged in a series according to their strentgh which is called Spectrochemical Series. weak Ligand

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